Chapter 6

1. http://video.google.com/videoplay?docid=-2134266654801392897&ei=r5JES9WfBofCrQK59oi1Aw&q=cesium+in+water&hl=en&client=firefox-a#http://video.google.com/videoplay?docid=-2134266654801392897&ei=r5JES9WfBofCrQK59oi1Aw&q=cesium+in+water&hl=en&client=firefox-a# Chapter 6 The Periodic Table

2. Organizing the Elements • Chemists use properties of elements in order to sort them into groups. • J.W. Dobereiner (1780-1849) used triads • Early in the 19th cent., a number of chemists had noticed certain relationships between the properties of elements and their atomic weight. In 1829 J. W. Döbereiner stated that there existed some three-element groups, or triads, in which the atomic weight of the middle element was the average of the other two and the properties of this element lay between those of the other two. • For example… Li = 6.941 Na = 23.98977 average = 23.02 K = 39.0983

3. Mendeleev’s Periodic Table • Dmitri Mendeleev, in 1869, published a table arranged in order of increasing atomic mass • He also took into account ability to react with oxygen • Gaps in his model accommodated elements which had not yet been discovered.

4. Periodic Law • (1913-14) Henry Moseley did X-ray experiments that showed that the elements have different numbers of protons in the nucleus (AKA Atomic Number) • When elements are arranged in order of increasing atomic number, there is a periodic repetition of the physical and chemical properties

5. The Modern Periodic Law – the physical and chemical properties of the elements are periodic functions of their atomic number. It is important to remember that before our "modern" periodic law, the Periodic Table of Elements was organized according to increasing atomic weights, and not by atomic number. English chemist John Newlands and Russian chemist, Dmitri Mendeleev ("father of the Periodic Table"), both categorized the elements on the basis of increasing atomic weights. The present law (and method of organization) was applied first by Henry Moseley in 1914.

6. Metals, Nonmetals, & Metalloids • Metals (in general 3 or less valence electrons) • Most elements are metals ~ 80% • Good conductors of heat and electricity • Freshly cut metal will have a high luster, or sheen, caused by a metal’s ability to reflect light • Solid at room temperature (except Hg) • Ductile, malleable, and tenacious (not easily pulled apart)

7. Nonmetals (gen. 4+ valence electrons) • Solid, liquid, and gaseous nonmetals exist at room temperature (most are gases) • Poor conductors of heat and electricity • Brittle

8. www.elementsdatabase.com • Metalloids • In general behave like nonmetals, however they can conduct electricity • Their behavior can be controlled with changing conditions, ex. Silicon can become a good conductor or electricity when a small amount of boron is added • 7 elements bordering the zigzag line on the periodic table

9. Classifying the Elements • Group I - alkali metals • The Arabic al qalīy means “the ashes”, wood ash is rich in compounds of sodium and potassium Group 2 – alkali earth metals Group 17 – halogens • The Greek “hals”, means salt, and the Latin genesis, means “to be born” Group 18 – Noble Gasses

10. Properties of the groups • Group 1 (s-block)– sodium family – highly active, soft, grey, metals, one s electron • Group 2 (s-block)– calcium family – active metals, but not as active as the alkali metals- 2 s electrons • Group 13(p-block)– boron family – varies from non-metallic to metallic • Group 14 (p-block)– carbon family – nonmetal to metalloid to metal

11. Group 15 (p-block)– nitrogen family – same as carbon family but As and Sb have both characteristics • Group 16 (p-block) - oxygen family (sometimes called chalcogens) – nonmetals to radioactive metals • Group 17 (p-block)– Halogens – very active nonmetals • Group 18 (p-block)– Noble gases – only Kr, Xe, Rn, have had compound prepared

12. Electron Configurations in Groups • The Noble Gases - elements having a full outer shell of 8 valence e- (p6) • The Representative Elements – see groups mentioned earlier • AKA Main Group Elements • Placed together because the outer e- configuration gives these elements similar properties

13. Transition Elements • Group B Elements Transition metals – the highest occupied s sublevel and a nearby d sublevel contain electrons Inner transition metal – the highest occupied s sublevel and a nearby f sublevel generally contain elections • Lanthanide and Actinide series • Below the main block on the periodic table

14. Periodic Trends • Atomic Size • Atomic radius – one half the distance between the nuclei of two atoms of the same element when the atoms are joined (draw here) • Group trends - the atomic radius increases as we move down to a given group because of the addition of another energy level • Period trends - the atomic radius decreases as we move across a given period because of increased nuclear charge from the addition of protons

16. 11p+ 12n° 11p+ 12n° 10p+ 10n° Na has 11 electrons Na+ has 10 electrons Ne has 10 electrons Radius increases because shells are added Increased radius will make it easier to lose an electron because of greater distance between positive and negative charges (Remember the Li, Na, K Demo?). We’ll talk more about this later

17. Atomic # increases. More protons means greater attraction between nucleus and outer electron thus higher ionization energy. The greater attraction also means that outer electrons are brought closer to the nucleus, thus smaller atomic radius results. Li (ve= 1) Be (ve= 2) B (ve= 3) + + + + + + + + + + + +

18. You know what to do! Which of the follow has the largest atomic radius: Li, O, C, or F? Na, Li, K, or Rb

19. Ions + - • An ion is an atom or group of atoms that has a positive or negative charge • Cation – an ion with a positive charge (loss of valence electron) usually metals • Anion – an ion with a negative charge (gain of valence electron) usually nonmetals

20. Trends in Ionization Energy • Defined as the energy required to remove the most loosely-held electron from an atom (measure in kJ) • First ionization energy – energy needed to remove the first electron • Ionization energy decreases down a group and increases across a period

21. Group trends – IE generally decreases in a family • Periodic trends – IE generally increases in a period The general formula is… Ao + energy  A+ + electron

22. Arrange the following in order of least to greatest ionization energy… Li Be Na Answer… Be > Li > Na As more e- are removed, the IE greatly increases because of the greater positive charge (and hold) of the nucleus

23. Trends in Ionic Size • Metallic Ions: are smaller than the corresponding atom • electrons are lost so they are pulled on harder by the nucleus • there are fewer energy levels once the e-‘s are lost

24. Nonmetallic Ions: are larger than the corresponding atom • Since electrons are gained they are repelled more by the nucleus • Electrons outnumber protons so they are farther from the nucleus and the atom swells • Cations are always smaller than the atoms from which they are formed • Anions are always larger than the atoms from which they are formed

25. Explain in your own words why cations are always smaller and anions are always larger than their corresponding atoms.

26. Electronegativity A measure of the ability of an atom in a compound to attract electrons (to itself) from another atom in the compound Only happens with bonded atoms Electronegativity increases as you move across a period Electronegativity decreases as you move down a group (why?)

27. To Review…